Chapter 5

Chemical Bonding and Nomenclature

Objectives

• Describe covalent bonding

• Explain the energy changes as molecules are formed

Bonds

• Forces that hold atoms together and make them function as a unit

• Two types of bonds– Ionic – Transfer of electrons

• Between metals and nonmetals–Called Salts

– Covalent – Sharing of electrons• Between nonmentals

–Called Molecules

Molecules

• A group of nonmetallic atoms covalently bonded

• Can be different atoms such as H2O• Can be the same atom such as O2

– Diatomic element• H2, N2, O2, F2, Cl2, Br2, I2

• Chemical Formula – Tells the number and types of atoms present– CH4 has 1 Carbon and 4 Hydrogens

Teacher where do molecules come from?

• Well when two atoms are attracted to each other they come together to share and share a special type of bond. Once this bond is established a molecule is born.

• Seriously!

The Making a Molecule of H2

1 A is 1x10-10m

So . . .

10A is 1x10-9m

The Atoms Need a Little Push

They Are Attracted!

What Is Going to Happen?

A Molecule Is Formed!

Bond Length

• Distance between two bonded nuclei

• Bond length changes when different types of atoms are bonded together

• Long bonds tend to be weaker

• Short bonds tend to be stronger

A Molecule Is Formed!

Bond Length

Atomic Interactions

• There are electrostatic interactions that take place within the atom and cause bonds to form

• Proton/Proton

• Electron/Electron

• Proton/Electron

Bonding Forces

Bond Length Diagram

Energy Changes in Reactions

• Breaking Bonds requires an input of energy

– Endothermic

• Forming Bonds releases energy

– Exothermic

• How does this compare to Biology?

How Many Bonds Need To Form

• Enough bonds to give eight electrons

• Octet Rule – Works most of the time – Duet for Hydrogen (2)– Quartet for Beryllium (4)– Sextet for Boron (6)

Specific Example for Hydrogen

Fluorine

Hydrogen Chloride

Homework

• p.201 #37,40,41,43,45,46,48,53

Objectives

• Draw dot diagrams for atoms

• Predict the formulas of compounds

• Draw lewis structures for molecules and ions

Dot Diagrams

• A method for showing the number of valence electrons around an atom

• Place elements symbol in the middle and electrons are represented as dots on the top, bottom, left and right

• Example

Li

Valence Electrons by Group

12 4 5 67

83

• Examples – Draw Dot Diagrams for the following elements Na, Mg, Al, Si, P, S, Cl, Ne

MgNa

Si

Al

PS Cl Ne

Formula Prediction

• Why is water always H2O?

• We can use dot diagrams to predict formulas

• Make formulas that allow the compounds to share electrons so each atom has 8– Or 2 if H, 4 if Be, and 6 is B

• Example – Predict the formula of a compound that is made of only Hydrogen and Oxygen

• Example – Predict the formula of a compound that is made of only Silicon and Hydrogen

• Example – Predict the formula of a compound that is made of only Arsenic and Chlorine

Usual Number of Bonds Formed

14 3 21

Problems

• p201 #56,62,63,64

Lewis Structures

• A method for determining the arrangement of bonds in covalent species

• Similar to dot structures but shows all bonds present

How 2 Draw

• Determine the number of valence electrons

• Determine the central atom. – Usually the single atom, or the one in the middle of

the formula

• Place other atoms around the middle and bond

• Complete octets with remaining electrons

• If each atom does not have 8 electrons – Multiple Bonds my be necessary

Bond Types

• Single Bonds = 2 electrons– Weakest and longest bonds

• Double Bonds = 4 electrons– In the middle

• Triple Bonds = 6 electrons– Strongest and shortest

There are not quadruple bonds!

General Rules

• Hydrogen will only form single bonds

• Halogens usually only form 1 bond. Why?– 7 valence electrons

• Oxygen will have 2 bonds and often forms multiple bonds

• Carbon likes to form chains

• Example – Draw the Lewis Structure for CBr4

• Example – Draw the Lewis Structure for NF3

• Example – Draw the Lewis Structure for O2

• Example – Draw the Lewis Structure for CS2

• Example – Draw the Lewis Structure for BeF2

• Example – Draw the Lewis Structure for C6H14

Polyatomic Ions

• Atoms that are covalently bonded together and have a charge

• Lewis structure rules– Negatively charged add electrons– Positively charged subtract electrons– Place Lewis structure in brackets when you

are finished

• Example – Draw the Lewis Structure for NO+

Resonance

• Species where equivalent Lewis structures exist

• Electron density is spread out evenly between resonant bonds– Delocalized – Spread out

• Often present in polyatomic ions

• Example – Draw the Lewis Structure for CO3

-2

• Example – Draw the Lewis Structure for NO2

-

• Example – Draw the Lewis Structure for AsF5

Homework

• p202 #71,74,75,76,77,82

• #71 C2H4

• #74 SO2

• #75 O3

• #76 Any Similarities b/t 74 and 75

• #77 Justify that that bonds are 1.5 of a bond

• #82 NO3-

Objectives

• Explain ionic bonding

• Name ionic compounds, acids, and covalent compounds

• Write formulas for ionic compounds, acids and covalent compounds

Ionic Bonds

• Involves the transfer of electrons between atoms to form ions

• Usually takes place between metals and nonmetals

• The ions are oppositely charged and strongly attract each other

• Large numbers of ions come together to form a crystal of the ionic compound

Electron Transfer

Na Cl

Properties of Ionic Compounds

• High melting points

– Solids at room temperature

• Hard

• Brittle

• Do not conduct current as solids

• Conduct current as liquids, gas, solutions

• Why?

Properties of Ionic Compounds

• High melting points

– Solids at room temperature

• Hard

• Brittle

• Do not conduct current as solids

• Conduct current as liquids, gas, solutions

• Why?

Ion Formation

• Atoms gain or lose electrons to achieve the same number of electrons as noble gases

• Metals lose electrons – Become positive ions (cations)

• Nonmetals gain electrons– Become negative ions (anions)

• Many groups form specific ions of specific charge

Ion Charge by Group+1

+2 -3 -2 -1+3

Ionic Compounds

• Total charge on the compound is zero

• Formulas are reduced to lowest ratios

• Formulas indicate the ratio of ions present

Naming Ionic Compounds

• Cation is named first, anion second• Cation keeps its name, anion ends in –ide

– Ex. Sulfide, Chloride, Oxide

• CaCl2 – Calcium Chloride

• NaBr– Sodium Bromide

• Ba3P2

– Barium Phosphide

Naming Ionic Compounds II

• Some cations can have more than one charge– Transition metals, Tin, Lead

• A roman numeral is placed after the cation name to indicate charge

• FeI2

– Iron (II) Iodide

• FeI– Iron (I) Iodide

Cont.

• Exceptions to the transition metal rule– Silver, Zinc, Chromium only form one ion

• Silver only +1• Zinc only +2• Chromium only +2

– Therefore roman numeral is not required• AgBr

– Silver Bromide not Silver (I) Bromide

Polyatomic Ions

• A covalently bonded group of atoms that possess a charge.

• Table 2.5 MUST be memorized!

• One more oxygen that normal Per- -ate

• Normal number of oxygens -ate

• One less oxygen than normal -ite

• Two less oxygens than normal Hypo- -ite

Polyatomic Ions p. 193

Naming Rules

• -ate to –ite one less oxygen

• Per- one more oxygen that –ate

• Hypo- one less oxygen that –ite

• Hydrogen added to ion increases charge by +1

• Phosphate PO4-3

• Hydrogen Phosphate HPO4-2

• Dihydrogen Phosphate H2PO4-1

Naming Ionic Compounds III

• Name cation as normal – Include roman number if necessary

• Name polyatomic ion

• NaNO3

– Sodium Nitrate

• NiSO4

– Nickel (II) Sulfate

Naming Acids

• Two types of acids1. Hydrogen and halogen (Hyrdohalic Acids)

– Named Hydroelementic Acid– HCl Hydrochloric Acid

2. Hydrogen and polyatomic ions– Name depends on name of polyatomic

ion• -ate becomes –ic• -ite becomes – ous

Cont.

• Acid formulas with start with H

• H2SO4

– Sulfuric Acid

• HNO3

– Nitric Acid

• H2SO3

– Sulfurous Acid

• HNO2

– Nitrous Acid

Covalent Compounds

• Only have nonmetals

• Use Number prefixes

• 1 = Mono 6 = Hexa

• 2 = Di 7 = Hepta

• 3= Tri 8 = Octa

• 4 = Tetra 9 = Nona

• 5 = Penta 10= Deca

Cont.

• The first element gets its name– If something other than one use a prefix

• The second element ends in –ide– Use prefix no matter what

• CO– Carbon Monoxide

• CO2

– Carbon Dioxide

Writing Formulas from Names

• Work backwards from the name

– Determine charges of the ions

– Make ratio so charges balance

• Reduce formulas to lowest ratio

• Except formulas that are fixed

– Peroxide O2-2

Examples

• Calcium Fluoride– CaF2

• Calcium Nitrate– Ca(NO3)2

• Iron (III) Sulfate– Fe2(SO4)3

• Titanium (IV) Oxide– TiO2

Cont.

• Chloric Acid– HClO3

• Hydroiodic acid– HI

• Carbon Tetrabromide– CBr4

• Diphosphorus Tetrafluoride– P2F4

Odd Compounds

• Cobalt (IV) Oxide

– CoO2

– Just being reduced

– Peroxides tend to form w/ Group 1

• Mercury (I) Bromide

– Hg2Br2

– Mercury (I) must be Hg2+2

Homework

• p203 #’s 92,94,96,97, 121, 122, 124 new only,125,126,127,135